5.5.2 Diamond

This topic explains diamond — one of the most important allotropes of carbon — including its structure, properties, and uses.

Definition

Diamond is an allotrope of carbon in which the carbon atoms are arranged in a diamond cubic crystal lattice.

Simple explanation: Diamond is the hard, transparent form of carbon used in jewellery and cutting tools.

Main Idea

Thanks to the presence of strong covalent bonds and a rigid tetrahedral structure, diamond is the hardest material ever discovered.

Simple explanation: Diamond is hard because every carbon atom is locked into a rigid 3D network by strong covalent bonds.

Core Concept

Structure of Diamond

• Carbon atoms are arranged tetrahedrally.
• Each carbon atom is attached to four other carbon atoms.
• Distance between carbon atoms: 1.544 × 10⁻¹⁰ meter.
• C–C–C bond angle: 109.5°.
• It is a strong, rigid three-dimensional structure.
• This results in an infinite network of atoms.

Simple explanation: Think of diamond as a giant 3D puzzle where each carbon atom connects to four others in a perfect tetrahedral shape.

Key Properties Arising from Structure

Detailed Explanation

Structure of Diamond

• The carbon atoms are arranged tetrahedrally.
• Each carbon atom is attached to four other carbon atoms.
• The distance between carbon atoms is 1.544 × 10⁻¹⁰ meter.
• The C–C–C bond angle is 109.5°.
• It is a strong, rigid three-dimensional structure.
• This results in an infinite network of atoms.

• Every carbon atom is locked into position by four strong covalent bonds.
• There are no weak points in the structure.
• To scratch or break diamond, you would need to break these strong covalent bonds — which requires enormous energy.
• This accounts for diamond's hardness, extraordinary strength and durability.

• Diamond has a more closely packed structure (3.514 g/cm³) than graphite.
• The atoms are held tightly in a 3D network.

Simple explanation: Diamond's rigid 3D network makes it hard, strong, and dense.

Uses of Diamond

The book lists the following uses:

1. Ornaments: Diamonds are most commonly used in ornaments like rings, necklaces, earrings, etc. In the gem industry, the value of diamonds is very high. They are used in making jewellery because of their durability and lustre property.

2. Cutting and drilling: Its property of hardness is useful to drill, grind or cut materials. Hence, some blades used for cutting and drills in the industry use diamonds. They are present on the edges and tips in small sizes.

3. Medicines and beauty products: Diamonds are used in making medicines and beauty products. They are also used in making medical tools like tools used in cataract surgery. Nano-diamonds have potential health benefits.

4. Sound equipment: Diamonds produce high-quality sound because they are hard and vibrate easily at high speed. They are also used in DJ equipment and high-quality recorders.

Simple explanation: Diamond is used in jewellery, cutting tools, medical tools, and high-quality sound equipment.

Comparison with Graphite

Why Diamond Does Not Conduct Electricity

• Each carbon atom in diamond uses all 4 valence electrons in covalent bonds.
• There are no free (delocalized) electrons.
• So diamond is a non-conductor of electricity.

Simple explanation: In diamond, all electrons are locked in bonds — none are free to carry current.

Why Graphite Conducts but Diamond Does Not

• Graphite: Each C bonds to only 3 other C atoms → the 4th electron is delocalized → conducts electricity.
• Diamond: Each C bonds to 4 other C atoms → all electrons are used in bonds → no free electrons → does not conduct electricity.

Simple explanation: Graphite has a spare electron per carbon; diamond doesn't.

Intermediate Level

Why Diamond Is the Hardest Substance

• The structure is a giant covalent network.
• Every carbon atom is bonded to four others at a 109.5° angle.
• These bonds are strong covalent bonds.
• There are no weak planes or layers.
• To deform diamond, you would need to break strong covalent bonds — this requires enormous energy.
• So diamond is extremely hard.

Simple explanation: Diamond has no weak spots — every atom is locked in place by strong covalent bonds.

Why Diamond Has a High Melting Point

• Diamond has strong covalent bonds throughout its structure.
• Melting would require breaking these strong covalent bonds.
• This requires an enormous amount of energy.
• So diamond has a very high melting point (about 3550 °C).

Why Diamond Is a Non-Conductor

• Each carbon atom has 4 valence electrons.
• All 4 are used in covalent bonds with neighboring carbon atoms.
• There are no free (delocalized) electrons.
• So diamond cannot conduct electricity.

Uses in Detail

Important Points

• - Diamond: allotrope of carbon; atoms arranged in a diamond cubic crystal lattice.
• - Structure: each C bonded to 4 other C atoms; tetrahedral arrangement; C–C–C bond angle = 109.5°; bond distance = 1.544 × 10⁻¹⁰ m.
• - Properties:
• - Hardest material ever discovered
• - Extraordinary strength and durability
• - Higher density than graphite (3.514 g/cm³)
• - Non-conductor of electricity (no free electrons)
• - Very high melting point
• - Uses: jewellery, cutting/drilling tools, medical tools (cataract surgery), beauty products, sound equipment (DJ equipment, recorders).
• - Comparison with graphite: diamond is hard, non-conductive, 3D network; graphite is soft, conductive, 2D layers.

Remember

• - Diamond = allotrope of carbon.
• - Each C bonded to 4 others → 3D tetrahedral network.
• - Hardest material.
• - Non-conductor (no free electrons).
• - High density (3.514 g/cm³).
• - Uses: jewellery, cutting tools, medical tools, sound equipment.
• - Comparison: diamond (hard, non-conductor) vs graphite (soft, conductor).

Common Mistakes

• 1. Thinking diamond and graphite are different elements. They are both carbon — just different allotropes.
• 2. Thinking diamond conducts electricity. It does not — no free electrons.
• 3. Forgetting that diamond is the hardest substance. This is a key property.
• 4. Thinking diamond is only for jewellery. It has many industrial and medical uses.
• 5. Confusing diamond's structure with graphite's. Diamond = 3D tetrahedral; graphite = 2D layers.
• 6. Thinking all carbon allotropes have the same properties. They have very different structures and properties.

Quick Revision

• Diamond: allotrope of carbon; atoms arranged in a diamond cubic crystal lattice.
• Structure: each carbon atom is attached to four other carbon atoms at a 109.5° angle; bond distance 1.544 × 10⁻¹⁰ m; strong, rigid 3D tetrahedral network.
• Properties:
• Hardest material ever discovered
• Extraordinary strength and durability
• Higher density than graphite (3.514 g/cm³)
• Non-conductor of electricity (no free electrons)
• Uses:
• Ornaments (rings, necklaces, earrings)
• Cutting and drilling tools
• Medical tools (cataract surgery)
• Medicines and beauty products
• DJ equipment and high-quality recorders

60-Second Revision

• Diamond = allotrope of carbon.
• Each C bonded to 4 others → 3D network.
• Hardest material.
• Non-conductor (no free electrons).
• Uses: jewellery, cutting tools, medical tools, sound equipment.
• Warning: diamond and graphite are both carbon — but have very different properties.